Metals and Non-metals

Elements can be classified as metals or non-metals on the basis of their properties.

Physical Properties of Metals

Examples of metals - iron, copper, aluminium, magnesium, sodium, lead, zinc.

Metals, in their pure state, have a shining surface. This property is called metallic lustre.
Metals are generally hard. The hardness varies from metal to metal.
Some metals can be beaten into thin sheets. This property is called malleability.
The ability of metals to be drawn into thin wires is called ductility. Gold is the most ductile metal.
Metals are good conductors of heat and have high melting points. The best conductors of heat are silver and copper. Lead and mercury are comparatively poor conductors of heat.
The metals that produce a sound on striking a hard surface are said to be sonorous.

All metals except mercury exist as solids at room temperature. Metals have high melting points but gallium and caesium have very low melting points. These two metals will melt if you keep them on your palm.

Alkali metals (lithium, sodium, potassium) are so soft that they can be cut with a knife. They have low densities and low melting points.

Physical Properties of Non-Metals

Some of the examples of non-metals are carbon, sulphur, iodine, oxygen, hydrogen, etc. The non-metals are either solids or gases except bromine which is a liquid.

Iodine is a non-metal but it is lustrous.

Carbon is a non-metal that can exist in different forms. Each form is called an allotrope. Diamond, an allotrope of carbon, is the hardest natural substance known and has a very high melting and boiling point. Graphite, another allotrope of carbon, is a conductor of electricity.

Chemical Properties of Metals

Examples: Aluminium, copper, iron, lead, magnesium, zinc and sodium.

Reaction of Metals with Oxygen

Almost all metals combine with oxygen to form metal oxides.

Metal + Oxygen → Metal oxide

For example, when copper is heated in air, it combines with oxygen to form copper(II) oxide, a black oxide.

$$ \ce{2Cu(s) + O2(g) -> 2CuO(s)} $$

Similarly, aluminium forms aluminium oxide.

$$ \ce{4Al(s) + 3O2(g) -> 2Al2O3(s)} $$

Magnesium burns in air with a dazzling white flame.

Metal oxides are basic in nature. But some metal oxides, such as aluminium oxide, zinc oxide show both acidic as well as basic behaviour. Such metal oxides which react with both acids as well as bases to produce salts and water are known as amphoteric oxides.

Aluminium oxide reacts in the following manner with acids and bases:

$$ \ce{Al2O3(s) + 6HCl(aq) -> 2AlCl3(aq) + 3H2O(l)} $$

$$ \ce{Al2O3 + 2NaOH -> 2NaAlO2 + H2O} $$

Most metal oxides are insoluble in water but some of these dissolve in water to form alkalis. Sodium oxide and potassium oxide dissolve in water to produce alkalis.

$$ \ce{Na2O(s) + H2O(l) -> 2NaOH(aq)} $$

$$ \ce{K2O(s) + H2O(l) -> 2KOH(aq)} $$

All metals do not react with oxygen at the same rate. Different metals show different reactivities towards oxygen. Metals such as potassium and sodium react so vigorously that they catch fire if kept in the open. Hence, to protect them and to prevent accidental fires, they are kept immersed in kerosene oil.

At ordinary temperature, the surfaces of metals such as magnesium, aluminium, zinc, lead, etc., are covered with a thin layer of oxide. The protective oxide layer prevents the metal from further oxidation. Iron does not burn on heating but iron filings burn vigorously when sprinkled in the flame of the burner. Copper does not burn, but the hot metal is coated with a black coloured layer of copper(II) oxide. Silver and gold do not react with oxygen even at high temperatures.

Reaction of Metals with Water

Metals react with water and produce a metal oxide and hydrogen gas. Metal oxides that are soluble in water dissolve in it to further form metal hydroxide. But all metals do not react with water.

Metal + Water → Metal oxide + Hydrogen

Metal oxide + Water → Metal hydroxide

Metals like potassium and sodium react violently with cold water. In case of sodium and potassium, the reaction is so violent and exothermic that the evolved hydrogen immediately catches fire.

$$ \ce{2K(s) + 2H2O(l) -> 2KOH(aq) + H2(g) + \text{heat energy}} $$

$$ \ce{2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g) + \text{heat energy}} $$

The reaction of calcium with water is less violent. The heat evolved is not sufficient for the hydrogen to catch fire. Calcium starts floating because the bubbles of hydrogen gas formed stick to the surface of the metal.

$$ \ce{Ca(s) + 2H2O(l) -> Ca(OH)2(aq) + H2(g)} $$

Magnesium does not react with cold water. It reacts with hot water to form magnesium hydroxide and hydrogen. It also starts floating due to the bubbles of hydrogen gas sticking to its surface.

Metals like aluminium, iron and zinc do not react either with cold or hot water. But they react with steam to form the metal oxide and hydrogen.

$$ \ce{2Al(s) + 3H2O(g) -> Al2O3(s) + 3H2(g)} $$

$$ \ce{3Fe(s) + 4H2O(g) -> Fe3O4(s) + 4H2(g)} $$

Metals such as lead, copper, silver and gold do not react with water at all.

Reaction of Metals with Acids

Metals react with acids to give a salt and hydrogen gas.

Metal + Dilute acid → Salt + Hydrogen

$$ \ce{Mg(s) + 2HCl(aq) -> MgCl2(aq) + H2(g)} $$

$$ \ce{2Al(s) + 6HCl(aq) -> 2AlCl3(aq) + 3H2(g)} $$

$$ \ce{Zn(s) + 2HCl(aq) -> ZnCl2(aq) + H2(g)} $$

$$ \ce{Fe(s) + 2HCl(aq) -> FeCl2(aq) + H2(g)} $$

Hydrogen gas is not evolved when a metal reacts with nitric acid. It is because HNO₃ is a strong oxidising agent. It oxidises the H₂ produced to water and itself gets reduced to any of the nitrogen oxides (N₂O, NO, NO₂). But magnesium (Mg) and manganese (Mn) react with very dilute HNO₃ to evolve H₂ gas.

The rate of formation of bubbles was the fastest in the case of magnesium. The reaction was also the most exothermic in this case. The reactivity decreases in the order Mg > Al > Zn > Fe. In the case of copper, no bubbles were seen and the temperature also remained unchanged. This shows that copper does not react with dilute HCl.

Reaction of Metals with Solutions of other Metal Salts

Reactive metals can displace less reactive metals from their compounds in solution or molten form.

Metal A + Salt solution of B → Salt solution of A + Metal B

The reactivity series is a list of metals arranged in the order of their decreasing activities.

K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au

Reaction of Metals and Non-metals

Noble gases, which have a completely filled valence shell, show little chemical activity. The reactivity of elements is tendency to attain a completely filled valence shell.

The compounds formed in this manner by the transfer of electrons from a metal to a non-metal are known as ionic compounds or electrovalent compounds. For example, NaCl, MgCl₂

Properties of Ionic Compounds

(i) Physical Nature

Ionic compounds are solids and are somewhat hard because of the strong force of attraction between the positive and negative ions. These compounds are generally brittle and break into pieces when pressure is applied.

(ii) Melting and Boiling Points

Ionic compounds have high melting and boiling points. This is because a considerable amount of energy is required to break the strong inter-ionic attraction.

(iii) Solubility

Electrovalent compounds are generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.

(iv) Conduction of Electricity

The conduction of electricity through a solution involves the movement of charged particles. A solution of an ionic compound in water contains ions, which move to the opposite electrodes when electricity is passed through the solution. Ionic compounds in the solid state do not conduct electricity because movement of ions in the solid is not possible due to their rigid structure. But ionic compounds conduct electricity in the molten state. This is possible in the molten state since the elecrostatic forces of attraction between the oppositely charged ions are overcome due to the heat. Thus, the ions move freely and conduct electricity.

Extraction of Metals

The earth’s crust is the major source of metals. Seawater also contains some soluble salts such as sodium chloride, magnesium chloride, etc. The elements or compounds, which occur naturally in the earth’s crust, are known as minerals.

At some places, minerals contain a very high percentage of a particular metal and the metal can be profitably extracted from it. These minerals are called ores.

Some metals are found in the earth’s crust in the free state. Some are found in the form of their compounds. The metals at the bottom of the activity series are the least reactive. They are often found in a free state. For example, gold, silver, platinum and copper are found in the free state. Copper and silver are also found in the combined state as their sulphide or oxide ores.

The metals at the top of the activity series (K, Na, Ca, Mg and Al) are so reactive that they are never found in nature as free elements. The metals in the middle of the activity series (Zn, Fe, Pb, etc.) are moderately reactive. They are found in the earth’s crust mainly as oxides, sulphides or carbonates.

The ores of many metals are oxides. This is because oxygen is a very reactive element and is very abundant on the earth.

Several steps are involved in the extraction of pure metal from ores.

Enrichment of Ores

Ores mined from the earth are usually contaminated with large amounts of impurities such as soil, sand, etc., called gangue. The impurities must be removed from the ore prior to the extraction of the metal.

The processes used for removing the gangue from the ore are based on the differences between the physical or chemical properties of the gangue and the ore. Different separation techniques are accordingly employed.

Extracting Metals Low in the Activity Series

Metals low in the activity series are very unreactive. The oxides of these metals can be reduced to metals by heating alone.

For example, cinnabar (HgS) is an ore of mercury. When it is heated in air, it is first converted into mercuric oxide (HgO). Mercuric oxide is then reduced to mercury on further heating.

$$ \ce{2HgS(s) + 3O2(g) ->[\text{Heat}] 2HgO(s) + 2SO2(g)} $$

$$ \ce{2HgO(s) ->[\text{Heat}] 2Hg(l) + O2(g)} $$

Similarly, copper which is found as Cu₂S in nature can be obtained from its ore by just heating in air.

$$ \ce{2Cu2S(s) + 3O2(g) ->[\text{Heat}] 2Cu2O(s) + 2SO2(g)} $$

$$ \ce{2Cu2O(s) + Cu2S(s) ->[\text{Heat}] 6Cu(s) + SO2(g)} $$

Extracting Metals Middle in the Activity Series

The metals in the middle of the activity series such as iron, zinc, lead, copper, are moderately reactive. These are usually present as sulphides or carbonates in nature. It is easier to obtain a metal from its oxide, as compared to its sulphides and carbonates. Therefore, prior to reduction, the metal sulphides and carbonates must be converted into metal oxides.

Roasting

The sulphide ores are converted into oxides by heating strongly in the presence of excess air. This process is known as roasting.

$$ \ce{2ZnS(s) + 3O2(g) ->[\text{Heat}] 2ZnO(s) + 2SO2(g)} $$

Calcination

The carbonate ores are changed into oxides by heating strongly in limited air. This process is known as calcination.

$$ \ce{ZnCO3(s) ->[\text{Heat}] ZnO(s) + CO2(g)} $$

The metal oxides are then reduced to the corresponding metals by using suitable reducing agents such as carbon. For example, when zinc oxide is heated with carbon, it is reduced to metallic zinc.

$$ \ce{ZnO(s) + C(s) -> Zn(s) + CO(g)} $$

Besides using carbon (coke) to reduce metal oxides to metals, sometimes displacement reactions can also be used. The highly reactive metals such as sodium, calcium, aluminium, are used as reducing agents because they can displace metals of lower reactivity from their compounds. For example, when manganese dioxide is heated with aluminium powder, the following reaction takes place

$$ \ce{3MnO2(s) + 4Al(s) -> 3Mn(l) + 2Al2O3(s) + Heat} $$

These displacement reactions are highly exothermic. The amount of heat evolved is so large that the metals are produced in the molten state.

The reaction of iron(III) oxide (Fe₂O₃) with aluminium is used to join railway tracks or cracked machine parts. This reaction is known as the thermit reaction.

$$ \ce{Fe2O3(s) + 2Al(s) -> 2Fe(l) + Al2O3(s) + Heat} $$

Extracting Metals Top in the Activity Series

The metals high up in the reactivity series are very reactive. They cannot be obtained from their compounds by heating with carbon. For example, carbon cannot reduce the oxides of sodium, magnesium, calcium, aluminium, etc., to the respective metals. This is because these metals have more affinity for oxygen than carbon.

These metals are obtained by electrolytic reduction. For example, sodium, magnesium and calcium are obtained by the electrolysis of their molten chlorides. The metals are deposited at the cathode (the negatively charged electrode), whereas, chlorine is liberated at the anode (the positively charged electrode).

Refining of Metals

The metals produced by various reduction processes are not very pure. They contain impurities, which must be removed to obtain pure metals. The most widely used method for refining impure metals is electrolytic refining.

Electrolytic Refining

Many metals, such as copper, zinc, tin, nickel, silver, gold, etc., are refined electrolytically. In this process, the impure metal is made the anode and a thin strip of pure metal is made the cathode. A solution of the metal salt is used as an electrolyte.

On passing the current through the electrolyte, the pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte is deposited on the cathode. The soluble impurities go into the solution, whereas, the insoluble impurities settle down at the bottom of the anode and are known as anode mud.

Corrosion

Silver articles become black after some time when exposed to air. This is because it reacts with sulphur in the air to form a coating of silver sulphide.

Copper reacts with moist carbon dioxide in the air and slowly loses its shiny brown surface and gains a green coat. This green substance is basic copper carbonate.

Iron when exposed to moist air for a long time acquires a coating of a brown flaky substance called rust.

Prevention of Corrosion

The rusting of iron can be prevented by painting, oiling, greasing, galvanising, chrome plating, anodising or making alloys.

Galvanisation is a method of protecting steel and iron from rusting by coating them with a thin layer of zinc.

Alloying is a very good method of improving the properties of a metal. For example, iron is the most widely used metal. But it is never used in its pure state. This is because pure iron is very soft and stretches easily when hot. But, if it is mixed with a small amount of carbon (about 0.05 %), it becomes hard and strong.

When iron is mixed with nickel and chromium, we get stainless steel, which is hard and does not rust.

An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. It is prepared by first melting the primary metal, and then, dissolving the other elements in it in definite proportions. It is then cooled to room temperature.

If one of the metals is mercury, then the alloy is known as an amalgam. The electrical conductivity and melting point of an alloy is less than that of pure metals. For example, brass, an alloy of copper and zinc (Cu and Zn), and bronze, an alloy of copper and tin (Cu and Sn), are not good conductors of electricity whereas copper is used for making electrical circuits. Solder, an alloy of lead and tin (Pb and Sn), has a low melting point and is used for welding electrical wires together.